Problem 22 - 43 Years JEE ADVANCED (1978-2020) + JEE MAIN Chapterwise & Topicwise Solved Papers Chemistry

Question

Molecular shapes of \(\mathrm{SF}_{4}, \mathrm{CF}_{4}\) and \(\mathrm{XeF}_{4}\) are (a) the same, with 2,0 and 1 lone pairs of electrons respectively (b) the same, with 1,1 and 1 lone pairs of electrons respectively (c) different, with 0,1 and 2 lone pairs of electrons respectively (d) different, with 1,0 and 2 lone pairs of electrons respectively

Step-by-Step Solution

Answer

(d) different, with 1, 0, and 2 lone pairs respectively.

Step 1: Determine the Lewis Structure for SF4

To start, we need the Lewis structure for \(\mathrm{SF}_{4}\). Sulfur has 6 valence electrons and fluorine has 7 valence electrons. Therefore, \(\mathrm{SF}_{4}\) has a total of 34 valence electrons (6 from sulfur + 28 from four fluorines). Distributing these electrons, sulfur bonds with four fluorine atoms and has two remaining electrons, resulting in one lone pair on sulfur.

Step 2: Determine Molecular Shape of SF4

\(\mathrm{SF}_{4}\) has 4 bonding pairs and 1 lone pair, which causes a seesaw shape due to the lone pair on the equatorial plane in a trigonal bipyramidal electronic geometry.

Step 3: Determine the Lewis Structure for CF4

Next, construct the Lewis structure for \(\mathrm{CF}_{4}\). Carbon has 4 valence electrons, and each fluorine has 7 valence electrons, totaling 32 valence electrons for \(\mathrm{CF}_{4}\). Carbon forms four bonds with fluorine atoms, using all its available electrons for bonding and resulting in no lone pairs on carbon.

Step 4: Determine Molecular Shape of CF4

\(\mathrm{CF}_{4}\) has 4 bonding pairs and 0 lone pairs, resulting in a tetrahedral shape as there are no lone pairs to distort it.

Step 5: Determine the Lewis Structure for XeF4

Now, consider \(\mathrm{XeF}_{4}\). Xenon has 8 valence electrons, and each fluorine has 7, for a total of 36 valence electrons. Xenon forms four bonds with fluorine and uses two pairs of its electrons as lone pairs.

Step 6: Determine Molecular Shape of XeF4

\(\mathrm{XeF}_{4}\) has 4 bonding pairs and 2 lone pairs. The electronic geometry is octahedral; however, the presence of two lone pairs makes the molecule adopt a square planar shape.

Step 7: Compare Molecular Shapes and Lone Pairs

Finally, compare the results: \(\mathrm{SF}_{4}\) has 1 lone pair and a seesaw shape, \(\mathrm{CF}_{4}\) has no lone pairs with a tetrahedral shape, and \(\mathrm{XeF}_{4}\) has 2 lone pairs with a square planar shape. This shows the molecular shapes are different: \(\mathrm{SF}_{4}\) and \(\mathrm{XeF}_{4}\) each having different numbers of lone pairs and geometries.

Key Concepts

Lewis structure

To determine molecular shapes, the Lewis structure is an essential starting point. A Lewis structure helps visualize how atoms connect in a molecule and shows the arrangement of electrons.
In any Lewis structure:

• Atoms are represented by their chemical symbols.
• Lines between atoms show bonds – each line stands for a shared pair of electrons.
• Dots around the atomic symbols represent lone pairs – non-bonded valence electrons.

Creating a Lewis structure involves counting the total number of valence electrons in the molecules. These electrons can be bonding or non-bonding. For \( \mathrm{SF}_4 \), the formation begins with the sulfur atom sharing its electrons with four fluorine atoms, leaving a lone pair. For \( \mathrm{CF}_4 \), carbon completely uses up its valence electrons without any lone pairs being left. \( \mathrm{XeF}_4 \) involves xenon, which bonds with four fluorines and retains two lone pairs.
Understanding this structure is critical to predicting further molecular shape.

Lone pairs

Lone pairs are valence electrons that do not participate in bonding. They are crucial in determining the molecular geometry because they occupy space around an atom. Although not visible in the molecule's "sticks" and "balls," lone pairs have an electrostatic impact on molecular shape.
In VSEPR theory (Valence Shell Electron Pair Repulsion theory):

• Lone pairs repel more strongly than bonding pairs, leading to adjustments in bond angles.
• These pairs influence the overall geometry due to their spatial properties.

For example, in \( \mathrm{SF}_4 \), the lone pair position causes the seesaw shape as it pressures bonded atoms to rearrange. In \( \mathrm{XeF}_4 \), the pair induces a square planar configuration.
Each lone pair affects the three-dimensional structure and directly impacts how the electrons distribute around the central atom.

Electronic geometry

Electronic geometry considers the spatial arrangement of all electron pairs, both bonding and lone pairs, around a molecule's central atom. This holistic view includes every region of negative electron density.
For instance, in \( \mathrm{SF}_4 \), the electronic geometry can be characterized as trigonal bipyramidal because it includes four bonds and one lone pair. However, when only shaping the visible molecular part, the lone pair modifies its structure into a seesaw form. \( \mathrm{CF}_4 \) maintains a tetrahedral electronic geometry, as there are no lone pairs to alter the symmetrical shape. Meanwhile, \( \mathrm{XeF}_4 \) features octahedral electronic geometry encompassing four bonds and two lone pairs.
Recognizing electronic geometry is key to predicting molecules' behaviors and interactions because it underpins why lone pairs result in their unique structures and orientations. By understanding these spatial and electronic influences, one can better understand molecular reactivity and properties.